Article pubs.acs.org/est
Atmospheric Fate of a Series of Carbonyl Nitrates: Photolysis Frequencies and OH-Oxidation Rate Constants R. Suarez-Bertoa,† B. Picquet-Varrault,*,† W. Tamas,† E. Pangui,† and J-F. Doussin†,‡ †
Laboratoire Interuniversitaire des Systèmes Atmosphériques (LISA), UMR-CNRS 7583, Université Paris-Est Créteil et Université Paris Diderot, Institut Pierre Simon Laplace (IPSL), 61 Avenue du Général de Gaulle - 94010 - Créteil, France ‡ University of Colorado, CIRES, Boulder, Colorado 80309, United States S Supporting Information *
ABSTRACT: Multifunctional organic nitrates are potential NOx reservoirs whose atmospheric chemistry is somewhat little known. They could play an important role in the spatial distribution of reactive nitrogen species and consequently in ozone formation and distribution in remote areas. In this work, the rate constants for the reaction with OH radical and the photolysis frequencies of α-nitrooxyacetone, 3-nitrooxy-2-butanone, and 3-methyl-3-nitrooxy-2-butanone have been determined at room temperature at 1000 mbar total pressure of synthetic air. The rate constants for the OH oxidation were measured using the relative rate technique, with methanol as reference compound. The following rate constants were obtained for the reaction with OH: kOH = (6.7 ± 2.5) × 10−13 cm3 molecule−1 s−1 for α-nitrooxyacetone, (10.6 ± 4.1) × 10−13 cm3 molecule−1 s−1 for 3-nitrooxy-2-butanone, and (2.6 ± 0.9) × 10−13 cm3 molecule−1 s−1 for 3-methyl-3nitrooxy-2-butanone. The corresponding photolysis frequencies extrapolated to typical atmospheric conditions for July first at noon at 40° latitude North were (4.8 ± 0.3) × 10−5 s−1, (5.7 ± 0.3) × 10−5 s−1, and (7.4 ± 0.2) × 10−5 s−1, respectively. The data show that photolysis is a major atmospheric sink for these organic nitrates.
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INTRODUCTION In recent decades, several field campaigns have reported a significant deficit in the reactive nitrogen species (NOy) budget. This so-called “NOy shortfall” can reach up to 45% of the total concentration observed for the major reactive nitrogen species in the atmosphere.1−5 These “missing species” seem to originate from photochemical processes and it has been proposed that organic nitrates, other than PAN, could contribute to the nitrogen shortfall.6 Moreover, Buhr et al. have pointed out that the scale/occurrence of missing NOy is correlated with the photochemical age of the considered air masses, suggesting that organic and multifunctional nitrates, formed by the oxidation of the parent hydrocarbons, may constitute an important fraction. Such species can be long-lived in the atmosphere acting as NOx reservoirs,2 which can be transported downwind to remote regions where their decomposition, and subsequent NOx release, will influence the regional ozone production.2,7−10 Organic nitrates are formed in NOx rich air by the degradation of VOCs through two main processes: • Rapid reaction with peroxy radicals, produced by the oxidation of VOCs: RO2 + NO → RO + NO2
(1a)
RO2 + NO → RONO2
(1b) © 2012 American Chemical Society
Reaction 1a is generally the main pathway. Nevertheless, reaction 1b becomes increasingly important with increasing peroxy radical carbon chain length.11 • Oxidation at the double bond of an unsaturated VOC by NO3 radical, which proceeds mainly by addition of the nitrate radical on the double bound to produce nitroalkyl radicals that can evolve into organic nitrates. Hence, alkyl nitrates but also polyfunctional organic nitrates (such as β- and δ-hydrooxynitrates, dinitrates, and carbonyl nitrates) can be formed. In particular, carbonyl nitrates are known to be formed by the NO3-induced oxidation of unsaturated hydrocarbons.12−16 While Atlas7 states that alkyl nitrates represent 20% of the NOy, others5,17 point out that in aged air masses, alkyl nitrates do not represent a large fraction of the missing NOy. Therefore, some other species must account for the missing NOy and model studies have suggested that multifunctional nitrates could form an important part of that fraction.18 Furthermore, field measurements show that alkyl nitrates, alone, do not account for the whole nitrogen shortfall observed.1 Field studies Received: Revised: Accepted: Published: 12502
July 20, 2012 October 31, 2012 November 5, 2012 November 5, 2012 dx.doi.org/10.1021/es302613x | Environ. Sci. Technol. 2012, 46, 12502−12509
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4000) which provide an actinic flux very comparable to that of the sun (see Figure 1). This similarity allows us to measure
also show that multifunctional nitrates can be present in polluted air masses.19,20 Beaver et al. have observed that first and second generation nitrates from isoprene along with a nitrate from MBO oxidation accounted for two-thirds of the alkyl nitrates observed over Sierra Nevada in the summer.21 α-Nitrooxyacetone represented an important fraction of the nitrates observed. Furthermore, Paulot et al. have suggested that α-nitrooxyacetone seems to be the longest-lived nitrate formed following the total oxidation of isoprene.22 Photolysis, reaction with OH radicals, thermal decomposition, and wet and dry depositions are considered to be the main removal processes of organic nitrates from the atmosphere. Of those, photolysis and reaction with OH radical are likely to be the most important sinks of some organic nitrates.23,24 In light of this, kinetic and mechanistic data associated with these removal pathways are crucial to evaluate their impact on the NOx budget. However, available data on organic nitrates removal processes, other than PAN and alkyl nitrates, is scarce. In fact, the only available information, to our knowledge, regarding the atmospheric chemistry of carbonyl nitrates is limited to the measurements of the absorption cross sections and photolysis rates of α-nitrooxyacetone, 1-nitrooxy-2butanone, and 3-nitrooxy-2-butanone,25 the Henry’s law coefficients of α-nitrooxyacetone,26 and finally the rate coefficients for the OH-oxidation of α-nitrooxyacetone, 1nitrooxy-2-butanone, and 3-nitrooxy-2-butanone.27 The present study aims at understanding the fate of carbonyl nitrates in the atmosphere. Hence, we provide rate constants for the photolysis and for the OH-oxidation reaction of three carbonyl nitrates: α-nitrooxyacetone, 3-nitrooxy-2-butanone, and 3methyl-3-nitrooxy-2-butanone.
Figure 1. Irradiation spectra comparison. The dashed curve is the solar spectrum calculated from TUV NCAR, 12:00 solar time, July 1 at 40° latitude North, overhead ozone column 200, and albedo 0.1. The solid curve is the spectra of xenon filtered by a 7-mm Pyrex film. The solar spectrum is scaled to the lamps spectrum to facilitate comparison of their respective shapes.
photolysis frequencies under realistic conditions. The chamber is mounted with a FT-IR spectrometer (Bruker Tensor 37) as well as NO, NO2, and O3 analyzers (Horiba). For more detailed information about this chamber the reader is referred to Wang et al.29 Concentrations of carbonyl nitrates were in the ppm range in a synthetic mixture of 80% N2 and 20% O2. Cyclohexane (2−4 ppm) was also added to the mixture as an OH-scavenger. At such a cyclohexane mixing ratio, it was estimated that more than 95% of the OH radicals were scavenged. All experiments were conducted at 298 ± 5 K. In order to assess the impact of the reactor’s walls and minimize their effects, the mixture was kept under dark conditions during 3 h after which it was irradiated during 4 h. Finally, it was left in the dark for an additional 30−60 min period (see Photolysis section below). Compounds were monitored by acquiring infrared spectra every 5 min (corresponding to 136 coadded interferograms) with a resolution of 0.5 cm−1 and a path length of 192 m. The photolysis rate of NO2 (JNO2) needed to be calculated in order to adjust the values obtained using the simulation chamber to atmospheric conditions corresponding to noon on July 1 at 40° latitude North (overhead ozone column 200, albedo 0.1) (see Photolysis section). Hence, 400 ppbv of NO2 in 1000 mbar of N2 was injected in CESAM and kept in the dark for 20 min. The lights were then turned on during 20 min, following which the mixture was left in the dark for an additional 20 min. The photolysis frequency was subsequently determined using a kinetic numeric model developed for previous NOx photo-oxidation experiments in CESAM.29 The fitting of modeled values from the measured data provided a NO2 photolysis frequency equal to 3.0 × 10−3 s−1. This value had an associated statistical error, 2σ, of 0.01. Chemicals and Gases. Dry synthetic air was generated using N2 (from liquid nitrogen evaporation, >99.995% pure, 99.995% pure, 99% Air Liquide) and NO2 (quality N20, >99% Air Liquide) were also used. Chemicals included hydroxyacetone (95% Alfa Aesar), 3hydroxy-2-butanone (97% Alfa Aesar), 3-hydroxy-3-methyl-2butanone (90% Alfa Aesar), cyclohexane (VWR), methanol (J.T. Baker), H2SO4 (95% VWR), NaNO2 (≥99 Prolabo), and isopropanol (VWR). Carbonyl nitrates 4−6 were synthesized using Kames’ method,30 a liquid/gas phase reaction in which the corresponding hydroxyketone (1−3) is reacted with NO3 radicals released from the dissociation of N2O5, at ice bath temperature, under dry conditions and stirring (Scheme 1). The carbonyl nitrates
wall
nitrate ⎯⎯⎯→ products k hν
nitrate → products J −
(2b)
d[nitrate] = (J + k)[nitrate] dt
Ln[nitrate]t = Ln[nitrate]0 − (k + J )t
By plotting Ln[nitrate]t vs. time, where [nitrate]t is the concentration of the carbonyl nitrate (nitrate) at time t, a straight line is obtained with a slope of (J + k). The same approach was applied to each of the “dark” periods, before and after irradiation, to determine their respective nitrate decay rates, namely kbefore and kafter. Once verified that kbefore and kafter were comparable, k was calculated as the average of kbefore and kafter for each experiment. However, for the first α-nitrooxyacetone experiment, only the kafter value was available and used to calculate J. Uncertainties. All the linear regressions for the kinetic experiments were performed using the method developed by Brauers and Finlayson-Pitts, which takes into account the uncertainties along both abscissa and ordinate axes. This method avoids systematic bias on the slope of the plot.34 For relative rate experiments, the overall uncertainties were calculated by adding the relative uncertainty corresponding to the statistical error and the error in the reference rate constant (here 20% for methanol). The statistical error was set at 2σ, (where σ is the standard deviation on the linear regression). In the case of the photolysis experiments, the uncertainties were calculated by adding the respective statistical errors (2σ) associated to the “dark” and “light” periods, the former set as the average of the uncertainties determined for both dark periods (vide inf ra), with the exception of the experiment 1 for both α-nitrooxyacetone and 3-nitrooxy-2-butanone. In the case of 3-nitrooxy-2-butanone, the uncertainty was not calculated using the approach detailed above, the statistical error being too low compared to the difference between kbefore and kafter. This difference was equal to 0.4 × 10−5 s−1 and, consequently an uncertainty of ±0.2 × 10−5 s−1 was used. In the case of αnitrooxyacetone, the highest uncertainty found for the complete pool of experiments (i.e., ± 0.2 × 10−5 s−1) was used to compensate for the missing kbefore. The overall uncertainty associated with the photolysis rate of each of the carbonyl nitrates was calculated as the average of the uncertainties obtained for each experiment, divided by the square root of the number of experiments.
Scheme 1
(4−6) and nitric acid were separated by liquid extraction using dichloromethane and water. The carbonyl nitrates’ structure and purity were verified by FT-IR and GC-MS. Only traces of other compounds were observed. Isopropyl nitrite was synthesized by dropwise addition of a dilute solution of H2SO4 into a mixture of NaNO2 and isopropanol.31 Determination of Rate Constants. The relative rate technique11 was used to determine the rate constants of the OH-induced oxidation of α-nitrooxyacetone, 3-nitrooxy-2butanone, and 3-methyl-3-nitrooxy-2-butanone, with methanol as a reference compound. We used the IUPAC recommended value k(methanol): 9.0 × 10−13 cm3 molecule−1 s−1 as the rate constant for the oxidation of methanol by OH radicals at 298 K and atmospheric pressure.32 We verified that photolysis and wall losses were negligible in the LISA simulation chamber for the studied carbonyl nitrates under our experimental conditions. Hence, it was assumed that reaction with OH was the only removal process for both the studied (nitrate) and the reference (methanol) compounds, and that neither reformed at any stage during the experiment. It can be shown33 that Ln
(2a)
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[nitrate]0 k [methanol]0 = nitrate Ln [nitrate]t k methanol [methanol]t
RESULTS AND DISCUSSION Oxidation by OH Radicals. Kinetic experiments were carried out, in triplicate, using the relative rate technique for each of the following carbonyl nitrates: α-nitrooxyacetone, 3nitrooxy-2-butanone, and 3-methyl-3-nitrooxy-2-butanone. Isopropyl nitrite was used as the OH source and methanol was used as the reference compound. The stability of the starting materials was verified over the time-scale of a typical experiment in different conditions and neither decomposition nor reactions were noticed. It was also confirmed that the carbonyl nitrates were stable when irradiated over the timescale of a typical experiment. Figure 2 represents Ln([nitrate]0/
where [nitrate]0 and [methanol]0, and [nitrate]t and [methanol]t stand for the concentration of the carbonyl nitrate and the reference compound at times 0 and t, respectively. The plot Ln([nitrate]0/[nitrate]t) vs Ln([methanol]0/[methanol]t) is linear with a slope equal to knitrate/kmethanol and an intercept of zero. Determination of Photolysis Frequencies. During the photolysis experiments, interaction/adsorption of the carbonyl nitrates with the stainless steel walls of CESAM were noticed. Therefore, reaction 2a was added to the system. Hence, once assumed that photolysis and wall interaction/adsorption remain 12504
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constant obtained for 3-nitrooxy-2-butanone ((10.6 ± 4.1) × 10 −13 cm 3 molecule −1 s −1 ) is also consistent, within uncertainties, with the value proposed in the same article. Moreover, 3-methyl-3-nitrooxy-2-butanone was found to be the least reactive of the three carbonyl nitrates studied, with a rate constant equal to (2.6 ± 0.9) × 10−13 cm3 molecule−1 s−1 (Table 1). It should be noted that, for this last compound, our study provides the first experimental value of the rate constant. The obtained rate constants were also compared with the values estimated using the structure−activity relationship (SAR) developed by Kwok and Atkinson35 to predict the rate constants for the OH-oxidation of VOCs (Table 1). To rationalize the difference in their respective reactivity, a closer look at the different molecules, their backbone as well as the functional groups attached, is needed (molecular structures are displayed in Table 1). According to Kwok and Atkinson SAR model the reactivity of the carbon 1 is the same in all three molecules. We observed that the reactivity of these molecules with OH radical increases with the substitution level of the position 3, until we get a nonreactive carbon (i.e., quaternary). Hence, the 3-nitrooxy-2-butanone, with a tertiary carbon in position 3, was found to be the most reactive, followed by the nitrooxyacetone with a secondary carbon in the same position, and, finally, the 3-methyl-3-nitrooxy-2-butanone where the carbon 3 is quaternary. However, according to the Kwok and Atkinson SAR model, and using the substituent factors for the nitro group (F(ONO2) = 0.14) from Talukdar et al.,36 the secondary carbon of the nitrooxyacetone is deactivated by the carbonyl group (F(=O) = 0.75) and even more by the nitro group (F(ONO2) = 0.14) and therefore should be the least reactive. That the 3-nitrooxy-2-butanone is more reactive than the nitrooxyacetone can be easily explained by the fact that position 3 corresponds to a tertiary carbon that remains more reactive than a secondary one deactivated by the same carbonyl and nitro groups (F(=O) = 0.75 and F(ONO2) = 0.14, respectively). Furthermore, it has one additional reactive carbon (carbon 4), a primary carbon, which, even though deactivated by the nitro substituent (F(>CHONO2) = 0.28), is strongly activated by the carbonyl group in β position(F(>CHC(O)) = 3.9). Finally, based on the SAR one would expect the 3-methyl-3-nitrooxy-2-butanone to be almost as reactive as 3-nitrooxy-2-butanone because the carbonyl activating effect (F(CC(O)) = 3.9) on the two primary carbons in position 4 and 5, should counterbalance the deactivation exerted by the nitro group (F(CONO2) = 0.28) and the loss of reactivity of position 3. However, our experimental determinations disagree with this prediction. The substituent factors proposed by Kwok and Atkinson35 on the one hand and Talukdar et al.36 on the other have been proven adequate to calculate reaction rate values for monofunctional ketones and monofunctional nitrates, respectively. However, we have shown in this work that they are not appropriate to estimate the reaction rates of carbonyl nitrates. Therefore, we attempted to calculate potentially more suitable F*(ONO2) and F*(CONO2) parameters using our experimental results. Hence, assuming that kCH3, kCH2, kCH, F(CH3), F(CH2), F(>CH), F(=O), and F( CC(O)) were correct, α-nitrooxyacetone was used to calculate the factor F*(ONO2), since only the carbon 3 is affected by the nitro group in α position, while 3-methyl-3nitrooxy-2-butanone was used to determine the factor F*( CONO2) since carbons 4 and 5 are both affected by the nitro
Figure 2. Kinetic plot for the reaction of OH radicals with 3-nitrooxy2-butanone (black solid symbols), α-nitrooxyacetone (black open symbols), and 3-methyl-3-nitrooxy-2-butanone (gray symbols). Within each series, the different symbols correspond to different experiments. The Ln([3-nitrooxy-2-butanone]0/[3-nitrooxy-2-butanone]t) have been Y shifted by 0.07 and the Ln([α-nitrooxyacetone]0/[ αnitrooxyacetone]t) have been Y shifted by 0.03.
[nitrate]t) vs Ln([methanol]0/[methanol]t) (see rate constant calculation above) for α-nitrooxyacetone, 3-nitrooxy-2-butanone, and 3-methyl-3-nitrooxy-2-butanone, respectively. It can be observed that all experiments are in good agreement for each carbonyl nitrate (Figure 2). Therefore, all experimental points were combined to provide the final rate constant for each compound. The obtained rate constants are given and compared to data from the literature in Table 1. The rate constant for the αnitrooxyacetone has been found equal to (6.7 ± 2.5) cm3 molecule−1 s−1. This value is consistent, within uncertainties, with the upper limit proposed by Zhu et al.,27 which is, to our knowledge, the only available experimental study in the literature on the OH-oxidation of carbonyl nitrates. The rate Table 1. OH Reaction Rate Constants of α-Nitrooxyacetone, 3-Nitrooxy-2-Butanone, and 3-Methyl-3-Nitrooxy-2Butanone; Comparison with Previous Data
*
n-Butane used as reference compound. 12505
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Table 2. Photolysis Rates of α-Nitrooxyacetone, 3-Nitrooxy-2-Butanone, and 3-Methyl-3-Nitrooxy-2-Butanone; Comparison with Previous Data compound α-nitrooxyacetone
3-nitrooxy-2-butanone
3-methyl-3-nitrooxy-2-butanone
experiment 1 2 average
1 2 average
1 2 average
kbeforea × 10−5
kafterb × 10−5
(k + J)c × 10−5
0.39 ± 0.05
0.64 ± 0.07 0.36 ± 0.05
2.12 ± 0.02 1.79 ± 0.02
0.59 ± 0.02 0.44 ± 0.01
0.14 ± 0.01 0.09 ± 0.01
0.20 ± 0.01 0.44 ± 0.01
0.10 ± 0.01 0.11 ± 0.01
2.10 ± 0.02 2.25 ± 0.02
2.48 ± 0.05 2.36 ± 0.05
Jd × 10−5 1.5 1.4 1.5 4.8
